The document provides a comprehensive overview of the fundamentals of chemistry, focusing on matter, elements, atoms, and atomic structure. It explains the characteristics and behaviors of elements, including atomic number, mass number, isotopes, and the formation of molecules and compounds. Additionally, it covers chemical bonding theories such as ionic and covalent bonding, as well as intermolecular forces like hydrogen bonding and van der Waals interactions.

1. The Foundations of Chemistry
Kanchanjunga Prasai, Ph.D.
Department of Biotechnology
School of Science
Kathmandu University
kjprasai@gmail.com

2. • The term matter refers to anything that occupies space and has mass.
• All matter is made up of substances called elements, which have
unique physical and chemical properties.
• An element is a pure or raw chemical substance that consists of only
one kind of atom and that cannot be further broken down by any
physical or chemical methods.
• An atom is defined as the smallest particle of an element.
• In nature, most elements exist as populations of atoms, either
separated or in contact with each other, depending on the
physical state.
Elements and Atoms

3. • Figure on the right depicts atoms of an
element in its gaseous state.
• Noble gases like helium, argon, etc., are
examples of monoatomic gases.
• This is a natural gold crystal. Gold sometimes
occurs in nature as a pure element.
• Each element has a name, such as gold, silicon,
oxygen, or copper; pure gold is composed of
only one type of atom, gold atoms.
Elements and Atoms

4. • The macroscopic properties of a piece of gold, such as color, density,
and combustibility, are different from those of a piece of copper
because the submicroscopic properties of gold atoms are different
from those of copper atoms.
• Thus, each element is unique because the properties of its atoms are
unique.
• Each element is designated by its chemical symbol, which is a single
capital letter or, when the first letter is already “taken” by another
element, a combination of two letters. For e.g., C for carbon and
Ca for calcium.
• Each element is identified according to the number of protons (atomic
number or Z) it has in its nucleus. In the modern periodic table, the
elements are listed in order of increasing atomic number.
Atom and Atomic Number

5. Mass Number and Isotopes
• The number of protons define the identity of an element (i.e., an
element with 6 protons is a carbon atom, no matter how many
neutrons may be present).
• The number of protons determines how many electrons surround the
nucleus, and it is the arrangement of these electrons that determines
most of the chemical behavior of an element.
• The mass number (A) is the total number of protons and neutrons in
the nucleus of an atom.
• Each proton and each neutron contributes one unit to the mass
number. Thus, a uranium atom with 92 protons and 146 neutrons in its
nucleus has a mass number of 238.
• Isotopes are members of a family of an element that all have the same
number of protons but different numbers of neutrons, and hence
differ in mass number as well as relative atomic mass.

6. Depiction of the atomic number, mass number, and symbol for four atoms,
two of which are isotopes of the element uranium (the nuclei are not drawn
to scale.)
Atomic Number, Mass Number, and Isotopes

7. Atoms and Subatomic Particles
• The first chemist to use the name ‘atom’ was John Dalton (1766-1844);
he used the word ‘atom’ to mean the smallest particle of an element.
• The protons and neutrons in an atom cluster together in the central
part, called the nucleus, and the electrons 'orbit' the nucleus.
• A particular atom will have the same number of protons and
electrons and most atoms have at least as many neutrons as protons.
• If the number of protons and electrons are equal, then the atom is
electrically neutral. If an atom has more or fewer electrons than
protons, then it has an overall negative or positive charge, respectively
– such atoms are called ions.
• Most atoms have a radius of about 10−10 m. So, the radii of atoms are
about 0.1 nm or 1 Å.

8. • The element hydrogen has the simplest
atoms, each with just one proton and
one electron. The proton forms the
nucleus, while the electron orbits
around it.
• All other elements have neutrons as
well as protons in their nucleus, such as
helium, which is shown on the right.
• The dot in the middle is the nucleus,
and the surrounding cloud represents
where the two electrons might be at any
time. The darker the shade, the more
likely that an electron will be there.
Atoms and Subatomic Particles
An illustration of the helium atom,
depicting the nucleus (pink) and the
electron cloud distribution (black).

9. • Atoms have a small positive nucleus surrounded by a much larger
region of space in which tiny negative electrons move continuously.
• The positive charge of the nucleus is due to positively charged
protons. The nucleus also contains uncharged neutrons which have
virtually same mass as protons.
• Protons and neutrons are about 1840 times heavier than electrons, so
virtually all the mass of an atom is concentrated in the nucleus.
Properties of Subatomic Particles

10. Molecules
• John Dalton went on to explain how atoms
could combine together to form molecules,
which he called ‘compound atoms’.
• A molecule is a particle containing two or
more atoms joined together chemically.
• Oxygen, for example, occurs in air as diatomic
(two-atom) molecules. Molecules of an element
• Oxygen consists of particles of O2 under ordinary conditions, but at
very high temperatures these split to form particles of O.
• So molecules of oxygen are written as O2 and atoms of oxygen are
written as O.

11. Differences Between Molecules and Compounds
Molecules Compounds
1. A molecule consists of two or
more atoms of the same or different
elements, that are chemically bound.
2. A molecule can contain two or
more atoms of the same element
joined together.
3. Therefore, all molecules are not
compounds.
4. For e.g., O2 is a molecule but not a
compound because it consists of
only one type of atom.
1. A compound is a molecule that
contains at least two different
elements joined chemically.
2. But a compound must always
have at least two atoms of different
elements.
3. But all compounds are molecules.
4. CO2 is a compound because it
consists of atoms of different
elements joined together.

12. Relative Atomic Masses – The 12C Scale
• Chemist originally measured atomic masses relative to hydrogen.
• Hydrogen was chosen initially because it had the smallest atoms and
those could be assigned a relative atomic mass of 1.
• Later, when scientists realized that one element could contain atoms of
different mass (isotopes), it became necessary to choose a single
isotope as a reference standard for relative atomic masses.
• In 1961, the isotope carbon-12 (12C) was chosen as the new standard,
because carbon is a solid which is much easier to store and transport
than hydrogen (which is gas).
• On the 12C scale, atoms of the isotope carbon-12 are assigned a
relative atomic mass of exactly 12.

13. Relative Atomic Masses – The 12C Scale
• The mass of carbon-12 atom is defined as exactly 12 atomic mass units
(amu, also known as daltons, Da); one atomic mass unit is 1/12th the
mass of a carbon-12 atom.
• One 12C atom has a mass of 12 daltons (12 Da, or 12 amu).
• Based on this standard, the 1H atom has a mass of 1.008 amu; in other
words, a 12C atom has almost 12 times the mass of an 1H atom.
• The relative atomic mass (also called atomic weight) is a weighted
average of all of the isotopes of that element, in which the mass of
each isotope is multiplied by the abundance of that particular
isotope.

14. Relative Atomic Masses – The 12C Scale
• For instance, it can be determined experimentally that neon consists of
three isotopes:
• neon-20 (with 10 protons and 10 neutrons in its nucleus) with a mass
of 19.992 amu and an abundance of 90.48%,
• neon-21 (with 10 protons and 11 neutrons) with a mass of 20.994 amu
and an abundance of 0.27%,
• and neon-22 (with 10 protons and 12 neutrons) with a mass of 21.991
amu and an abundance of 9.25%.
• The average atomic mass of neon is thus:
0.9048 × 19.992 amu = 18.09 amu
0.0027 × 20.994 amu = 0.057 amu
0.0925 × 21.991 amu = 2.03 amu
20.18 amu
The relative atomic mass
(atomic weight) is a weighted
average of all of the isotopes
of that element, in which the
mass of each isotope is
multiplied by the abundance
of that particular isotope.

15. Relative Atomic Masses – The 12C Scale
• The relative atomic mass of
carbon is 12.011. This means that
the average mass of a carbon
atom is 12.011, not 12.000.
• This is because naturally
occurring carbon contains a few
atoms of carbon-13 and
carbon-14 mixed with those of
carbon-12.
• The atomic mass is an average value; thus, while no individual
carbon atom has a mass of 12.011 amu, in the laboratory, we consider a
sample of carbon to consist of atoms with this average mass.

16. Relative Atomic Masses – The 12C Scale
• The atomic mass is useful in chemistry when it is paired with the mole
concept: the atomic mass of an element, measured in amu, is the
same as the mass in grams of one mole of an element.
• Thus, since the atomic mass of iron is 55.847 amu, one mole of iron
atoms would weigh 55.847 grams.
• The same concept can be extended to molecules. One formula unit of
sodium chloride (NaCl) would weigh 58.44 amu (22.98977 amu for
Na + 35.453 amu for Cl), so a mole of sodium chloride would weigh
58.44 grams.
• One molecule of water (H2O) would weigh 18.02 amu (2×1.008 amu
for H + 15.9994 amu for O), and a mole of water molecules would
weigh 18.02 grams.

17. Chemical Bonding and Intermolecular Forces
The Electronic Theory of Chemical Bonding
• When elements form compounds, they either lose, gain, or share
electrons so as to achieve stable electron configurations similar to
the next higher or lower noble gas in the periodic table.
• For e.g., the electron structure of Na is 2, 8, 1 and Mg is 2, 8, 2. They
can lose electrons from their outermost shell to form positively
charge ions (Na+ and Mg2+). These ions have an electron structure
like neon, the previous noble gas.
• Similarly, S and Cl have the electron structure of 2, 8, 6 and 2, 8, 7,
respectively. They can gain electrons to form negatively charged
ions (S2− and Cl−). These ions have an electron structure as the
next noble gas argon.

18. Chemical Bonding and Intermolecular Forces
• Silicon has a electron structure of 2, 8, 4.
• Elements like these do not usually form ions in their compounds,
but they do obtain an electron structure similar to noble gas by
sharing electrons rather than by losing or gaining electrons.
Transfer of Electrons-Ionic (Electrovalent) Bonding
• When metals like sodium and magnesium react with non-metals
like oxygen and chlorine, typical ionic compounds are formed.
• When the reactions occur, electrons are transferred from the metal
to non-metal until the outer electron shells of the resulting ions are
identical to those of noble gases.

19. Transfer of Electrons − Ionic (Electrovalent) Bonding
• The formation of compounds such as sodium chloride and calcium
fluoride involves a complete transfer of electrons. In these
compounds, ions are held together by electrostatic attraction in
ionic (electrovalent) bonds.

20. Ionic Interactions
• Ionic interactions involves attraction
or repulsion of a charged group on
one molecule with charged group on
the same or another molecule.
• Attractive forces occurs between
oppositely charged structures, and
repulsive forces occur between same
charged structures.
• Negatively charged groups, such as the carboxylate group
(–COO–) in the side chain of aspartate or glutamate, can interact
with positively charged groups, such as the amino group
(–NH3
+) in the side chain of lysine.

21. Sharing Electrons − Covalent Bonding
• This suggests that in molecules such as H2, electrons are shared by
two hydrogen atoms.
• Each hydrogen atom has only one electron.
• If the two hydrogen atoms come close together, their orbitals can
overlap.
• Figure on the right shows an
electron density map for a
hydrogen molecule.
• Although the highest concentration
of electrons is near each nucleus,
there is also a high concentration of
electrons between the two nuclei.

22. Sharing Electrons − Covalent Bonding
• The shared pair of electrons has resulted in a covalent bond.
• Therefore, a covalent bond involves the sharing of a pair of
electrons between two atoms. In a normal covalent bond, each
atom contributes one electron to the shared pair.
• The pair of electrons are then
attracted to each nucleus and shared
by each atom.
• Each hydrogen atom now has two
electrons, which is the same electronic
structure as helium and the H2
molecule is much more stable than
an H atom.

23. Co-ordinate (Dative Covalent) Bonding
• A co-ordinate bond (dative covalent bond) involves the sharing of
a pair of electrons between two atoms, both electrons in the bond
being donated by one atom.
• The word ‘dative’ meaning ‘giving’ is used because one atom gives
both the electrons in forming the bond.
• Once a co-ordinate bond has formed, it is very much like a normal
covalent bond.
• The arrow points from the atom donating the pair of electrons to
the atom accepting the electron pair.

24. Hydrogen Bonding
Polar Bonds
• When a covalent bond is between two identical atoms, such as O2,
the electrons are shared equally.
• But the electrons in a covalent bond will not be shared equally if
the atoms joined by the bond are different.
• The nucleus of one atom will attract the electrons in the bond
more strongly than that of the other atom.
• As a result of this, one end of the bond will have a small excess of
negative charge, δ− and the other end of the bond will have a small
deficit of negative charge resulting in a partial positive charge, δ+.

25. Hydrogen Bonding
Polar Bonds
• Covalent bonds in which this charge separation
occurs are described as polar and the charge
separation is known as bond polarity.
• Polar molecules, like HCl, have a tiny positive
electric pole, labelled δ+ and a tiny negative
electric pole labelled δ−; these two poles of
opposite charge in a molecule are described as a
dipole.
• The figure shows the polar covalent bond in a molecule of
hydrogen chloride.

26. Hydrogen Bonding
Electronegativity
• Electronegativity values determines the polarity of covalent
bonds; the greater the difference in electronegativity between the
elements forming a covalent bond, the more polar the bond.
• The electronegativity of an atom is the power of that atom in a
molecule to attract the shared electron in a covalent bond.
• Elements, such as fluorine, oxygen, nitrogen, chlorine, have higher
electronegativity values.
• These elements have stronger pull on shared electrons in a
covalent bond than elements such as hydrogen.
• So, in a polar O−−H bond, there is a δ− charge on the oxygen atom
and a δ+ charge on the hydrogen.

27. Hydrogen Bonding
• F, O, and N are three most electronegative elements.
• When they are bonded to an H atom, the electrons in the covalent
bond are drawn towards the electronegative atom.
• The H atom has no other electrons other than the one which it
shares in the covalent bond; those electrons are being pulled away
from it by the more electronegative N, O, or F.
• As the H atom has no inner shell of electrons, the single proton in
the nucleus is exceptionally ‘bare’ and very attractive to any lone
pair of electrons on another H2O, NH3, or HF molecule.
• Thus, H atoms attached to N, O, or F can interpose themselves
between two of these atoms.

28. Hydrogen Bonding
• The H atoms are covalently bonded to one of
the electronegative atoms and hydrogen
bonded to the other.
• The essential requirements for hydrogen
bonding are:
• a hydrogen atom bonded to a highly
electronegative atom, and
• an unshared pair of electrons on a second electronegative atom.
• Therefore, hydrogen bonding will usually occur in any substance
containing an H atom attached to an N, O, or F atom.
• How many hydrogen bonds can a single NH3 molecule form?

29. • The electrons in atoms and molecules are in continual motion.
• There will be more negative charge on the one side of the
molecule than on the other.
• The molecule has an instantaneous electric dipole and this dipole
will induce dipoles in neighboring molecules.
• At any particular moment, the electron
charge cloud around the molecule will
not be perfectly symmetrical.
van der Waals Forces (Interactions)

30. van der Waals Forces (Interactions)
• Negative dipoles will tend to induce positive dipoles and vice
versa; hence, the force between induced dipoles is always an
attraction.
• So the two dipoles weakly attract each other, bringing the two
nuclei closer.
• These weak attractions are called van der Waals interactions.

31. Hydrophobic Interactions
• Hydrophobic interactions describe the relations between water
and hydrophobes (low water-soluble molecules).
• Hydrophobic interactions result from the strong tendency of water
to exclude hydrophobic (nonpolar) groups or molecules.
• Alternatively, hydrophobic interaction refers to the tendency of
nonpolar compounds to self-associate in an aqueous environment.
Hydrophobic interactions between nonpolar molecules
No interaction Hydrophobic interaction

32. Hydrophobic Interactions
• These interactions arise because water
molecules prefer the stronger interactions
that they share with one another, compared
to their interaction with nonpolar molecules.
• It is these preferential interactions between
H2O molecules that “exclude” hydrophobic
substances from aqueous solution and drive
the tendency of nonpolar molecules to
cluster together.
• Thus, nonpolar regions of biological macromolecules are often
buried in the molecule’s interior to exclude them from the
aqueous milieu.

33. Four Types of Noncovalent (“Weak”) Interactions